Chapter 8 Concepts of Chemical Bonding

1
Chapter 8Concepts of Chemical Bonding
Chemistry, The Central Science, 10th
edition Theodore L. Brown H. Eugene LeMay, Jr.
and Bruce E. Bursten
John D. Bookstaver St. Charles Community
College St. Peters, MO ? 2006, Prentice Hall, Inc.
2
Chemical Bonds

• Three basic types of bonds
• Ionic
• Electrostatic attraction between ions
• Covalent
• Sharing of electrons
• Metallic
• Metal atoms bonded to several other atoms

3
Ionic Bonding
4
Energetics of Ionic Bonding

• As we saw in the last chapter, it takes 495
  kJ/mol to remove electrons from sodium.

5
Energetics of Ionic Bonding

• We get 349 kJ/mol back by giving electrons to
  chlorine.

6
Energetics of Ionic Bonding

• But these numbers dont explain why the reaction
  of sodium metal and chlorine gas to form sodium
  chloride is so exothermic!

7
Energetics of Ionic Bonding

• There must be a third piece to the puzzle.
• What is as yet unaccounted for is the
  electrostatic attraction between the newly formed
  sodium cation and chloride anion.

8
Lattice Energy

• This third piece of the puzzle is the lattice
  energy
• The energy required to completely separate a mole
  of a solid ionic compound into its gaseous ions.
• The energy associated with electrostatic
  interactions is governed by Coulombs law

9
Lattice Energy

• Lattice energy, then, increases with the charge
  on the ions.

• It also increases with decreasing size of ions.

10
Energetics of Ionic Bonding

• By accounting for all three energies (ionization
  energy, electron affinity, and lattice energy),
  we can get a good idea of the energetics involved
  in such a process.

11
Energetics of Ionic Bonding

• These phenomena also helps explain the octet
  rule.

• Metals, for instance, tend to stop losing
  electrons once they attain a noble gas
  configuration because energy would be expended
  that cannot be overcome by lattice energies.

12
Covalent Bonding

• In these bonds atoms share electrons.
• There are several electrostatic interactions in
  these bonds
• Attractions between electrons and nuclei
• Repulsions between electrons
• Repulsions between nuclei

13
Polar Covalent Bonds

• Although atoms often form compounds by sharing
  electrons, the electrons are not always shared
  equally.

• Fluorine pulls harder on the electrons it shares
  with hydrogen than hydrogen does.
• Therefore, the fluorine end of the molecule has
  more electron density than the hydrogen end.

14
Electronegativity

• The ability of atoms in a molecule to attract
  electrons to itself.
• On the periodic chart, electronegativity
  increases as you go
• from left to right across a row.
• from the bottom to the top of a column.

15
Polar Covalent Bonds

• When two atoms share electrons unequally, a bond
  dipole results.
• The dipole moment, ?, produced by two equal but
  opposite charges separated by a distance, r, is
  calculated
• ? Qr
• It is measured in debyes (D).

16
Polar Covalent Bonds

• The greater the difference in electronegativity,
  the more polar is the bond.

17
Lewis Structures

• Lewis structures are representations of
  molecules showing all electrons, bonding and
  nonbonding.

18
Writing Lewis Structures

• Find the sum of valence electrons of all atoms in
  the polyatomic ion or molecule.
• If it is an anion, add one electron for each
  negative charge.
• If it is a cation, subtract one electron for each
  positive charge.

• PCl3

5 3(7) 26
19
Writing Lewis Structures

1. The central atom is the least electronegative
   element that isnt hydrogen. Connect the outer
   atoms to it by single bonds.

Keep track of the electrons 26 ? 6 20
20
Writing Lewis Structures

1. Fill the octets of the outer atoms.

Keep track of the electrons 26 ? 6 20 ? 18 2
21
Writing Lewis Structures

1. Fill the octet of the central atom.

Keep track of the electrons 26 ? 6 20 ? 18
2 ? 2 0
22
Writing Lewis Structures

• If you run out of electrons before the central
  atom has an octet
• form multiple bonds until it does.

23
Writing Lewis Structures

• Then assign formal charges.
• For each atom, count the electrons in lone pairs
  and half the electrons it shares with other
  atoms.
• Subtract that from the number of valence
  electrons for that atom The difference is its
  formal charge.

24
Writing Lewis Structures

• The best Lewis structure
• is the one with the fewest charges.
• puts a negative charge on the most
  electronegative atom.

25
Resonance

• This is the Lewis structure we would draw for
  ozone, O3.

-
26
Resonance

• But this is at odds with the true, observed
  structure of ozone, in which
• both OO bonds are the same length.
• both outer oxygens have a charge of ?1/2.

27
Resonance

• One Lewis structure cannot accurately depict a
  molecule such as ozone.
• We use multiple structures, resonance structures,
  to describe the molecule.

28
Resonance

• Just as green is a synthesis of blue and yellow
• ozone is a synthesis of these two resonance
  structures.

29
Resonance

• In truth, the electrons that form the second CO
  bond in the double bonds below do not always sit
  between that C and that O, but rather can move
  among the two oxygens and the carbon.
• They are not localized, but rather are
  delocalized.

30
Resonance

• The organic compound benzene, C6H6, has two
  resonance structures.
• It is commonly depicted as a hexagon with a
  circle inside to signify the delocalized
  electrons in the ring.

31
Exceptions to the Octet Rule

• There are three types of ions or molecules that
  do not follow the octet rule
• Ions or molecules with an odd number of
  electrons.
• Ions or molecules with less than an octet.
• Ions or molecules with more than eight valence
  electrons (an expanded octet).

32
Odd Number of Electrons

• Though relatively rare and usually quite
  unstable and reactive, there are ions and
  molecules with an odd number of electrons.

33
Fewer Than Eight Electrons

• Consider BF3
• Giving boron a filled octet places a negative
  charge on the boron and a positive charge on
  fluorine.
• This would not be an accurate picture of the
  distribution of electrons in BF3.

34
Fewer Than Eight Electrons

• Therefore, structures that put a double bond
  between boron and fluorine are much less
  important than the one that leaves boron with
  only 6 valence electrons.

35
Fewer Than Eight Electrons

• The lesson is If filling the octet of the
  central atom results in a negative charge on the
  central atom and a positive charge on the more
  electronegative outer atom, dont fill the octet
  of the central atom.

36
More Than Eight Electrons

• The only way PCl5 can exist is if phosphorus has
  10 electrons around it.
• It is allowed to expand the octet of atoms on the
  3rd row or below.
• Presumably d orbitals in these atoms participate
  in bonding.

37
More Than Eight Electrons

• Even though we can draw a Lewis structure for the
  phosphate ion that has only 8 electrons around
  the central phosphorus, the better structure puts
  a double bond between the phosphorus and one of
  the oxygens.

38
More Than Eight Electrons

• This eliminates the charge on the phosphorus and
  the charge on one of the oxygens.
• The lesson is When the central atom is on the
  3rd row or below and expanding its octet
  eliminates some formal charges, do so.

39
Covalent Bond Strength

• Most simply, the strength of a bond is measured
  by determining how much energy is required to
  break the bond.
• This is the bond enthalpy.
• The bond enthalpy for a ClCl bond,
• D(ClCl), is measured to be 242 kJ/mol.

40
Average Bond Enthalpies

• This table lists the average bond enthalpies for
  many different types of bonds.
• Average bond enthalpies are positive, because
  bond breaking is an endothermic process.

41
Average Bond Enthalpies

• NOTE These are average bond enthalpies, not
  absolute bond enthalpies the CH bonds in
  methane, CH4, will be a bit different than the
• CH bond in chloroform, CHCl3.

42
Enthalpies of Reaction

• Yet another way to estimate ?H for a reaction is
  to compare the bond enthalpies of bonds broken to
  the bond enthalpies of the new bonds formed.

• In other words,
• ?Hrxn ?(bond enthalpies of bonds broken) ?
• ?(bond enthalpies of bonds formed)

43
Enthalpies of Reaction

• CH4(g) Cl2(g) ???
• CH3Cl(g) HCl(g)
• In this example, one
• CH bond and one
• ClCl bond are broken one CCl and one HCl bond
  are formed.

44
Enthalpies of Reaction

• So,
• ?Hrxn D(CH) D(ClCl) ? D(CCl) D(HCl)
• (413 kJ) (242 kJ) ? (328 kJ) (431 kJ)
• (655 kJ) ? (759 kJ)
• ?104 kJ

45
Bond Enthalpy and Bond Length

• We can also measure an average bond length for
  different bond types.
• As the number of bonds between two atoms
  increases, the bond length decreases.